The viscosity of a fluid is a measure of its resistance to deformation at a given rate. For liquids, it corresponds to the informal concept of "thickness": for example, syrup has a higher viscosity than water. Viscosity can be conceptualized as quantifying the frictional force that arises between adjacent layers of fluid that are in relative motion. For instance, when a fluid is forced through a tube, it flows more near the tube's axis than near its walls. In such a case, experiments show; this is because a force is required to overcome the friction between the layers of the fluid which are in relative motion: the strength of this force is proportional to the viscosity. A fluid that has no resistance to shear stress is known as an inviscid fluid. Zero viscosity is observed only at low temperatures in superfluids. Otherwise, the second law of thermodynamics requires all fluids to have positive viscosity. A fluid with a high viscosity, such as pitch, may appear to be a solid; the word "viscosity" is derived from the Latin "viscum", meaning mistletoe and a viscous glue made from mistletoe berries.
In materials science and engineering, one is interested in understanding the forces, or stresses, involved in the deformation of a material. For instance, if the material were a simple spring, the answer would be given by Hooke's law, which says that the force experienced by a spring is proportional to the distance displaced from equilibrium. Stresses which can be attributed to the deformation of a material from some rest state are called elastic stresses. In other materials, stresses are present which can be attributed to the rate of change of the deformation over time; these are called. For instance, in a fluid such as water the stresses which arise from shearing the fluid do not depend on the distance the fluid has been sheared. Viscosity is the material property which relates the viscous stresses in a material to the rate of change of a deformation. Although it applies to general flows, it is easy to visualize and define in a simple shearing flow, such as a planar Couette flow. In the Couette flow, a fluid is trapped between two infinitely large plates, one fixed and one in parallel motion at constant speed u.
If the speed of the top plate is low enough in steady state the fluid particles move parallel to it, their speed varies from 0 at the bottom to u at the top. Each layer of fluid moves faster than the one just below it, friction between them gives rise to a force resisting their relative motion. In particular, the fluid applies on the top plate a force in the direction opposite to its motion, an equal but opposite force on the bottom plate. An external force is therefore required in order to keep the top plate moving at constant speed. In many fluids, the flow velocity is observed to vary linearly from zero at the bottom to u at the top. Moreover, the magnitude F of the force acting on the top plate is found to be proportional to the speed u and the area A of each plate, inversely proportional to their separation y: F = μ A u y; the proportionality factor μ is the viscosity of the fluid, with units of Pa ⋅ s. The ratio u / y is called the rate of shear deformation or shear velocity, is the derivative of the fluid speed in the direction perpendicular to the plates.
If the velocity does not vary linearly with y the appropriate generalization is τ = μ ∂ u ∂ y, where τ = F / A, ∂ u / ∂ y is the local shear velocity. This expression is referred to as Newton's law of viscosity. In shearing flows with planar symmetry, it is what defines μ, it is a special case of the general definition of viscosity, which can be expressed in coordinate-free form. Use of the Greek letter mu for the viscosity is common among mechanical and chemical engineers, as well as physicists. However, the Greek letter eta is used by chemists and the IUPAC; the viscosity μ is sometimes referred to as the shear viscosity. However, at least one author discourages the use of this terminology, noting that μ can appear in nonshearing flows in addition to shearing flows. In general terms, the viscous stresses in a fluid are defined as those resulting from the relative velocity of different fluid particles; as such, the viscous stresses. If the velocity gradients are small to a first approximation the v
A molecule is an electrically neutral group of two or more atoms held together by chemical bonds. Molecules are distinguished from ions by their lack of electrical charge. However, in quantum physics, organic chemistry, biochemistry, the term molecule is used less also being applied to polyatomic ions. In the kinetic theory of gases, the term molecule is used for any gaseous particle regardless of its composition. According to this definition, noble gas atoms are considered molecules as they are monatomic molecules. A molecule may be homonuclear, that is, it consists of atoms of one chemical element, as with oxygen. Atoms and complexes connected by non-covalent interactions, such as hydrogen bonds or ionic bonds, are not considered single molecules. Molecules as components of matter are common in organic substances, they make up most of the oceans and atmosphere. However, the majority of familiar solid substances on Earth, including most of the minerals that make up the crust and core of the Earth, contain many chemical bonds, but are not made of identifiable molecules.
No typical molecule can be defined for ionic crystals and covalent crystals, although these are composed of repeating unit cells that extend either in a plane or three-dimensionally. The theme of repeated unit-cellular-structure holds for most condensed phases with metallic bonding, which means that solid metals are not made of molecules. In glasses, atoms may be held together by chemical bonds with no presence of any definable molecule, nor any of the regularity of repeating units that characterizes crystals; the science of molecules is called molecular chemistry or molecular physics, depending on whether the focus is on chemistry or physics. Molecular chemistry deals with the laws governing the interaction between molecules that results in the formation and breakage of chemical bonds, while molecular physics deals with the laws governing their structure and properties. In practice, this distinction is vague. In molecular sciences, a molecule consists of a stable system composed of two or more atoms.
Polyatomic ions may sometimes be usefully thought of as electrically charged molecules. The term unstable molecule is used for reactive species, i.e. short-lived assemblies of electrons and nuclei, such as radicals, molecular ions, Rydberg molecules, transition states, van der Waals complexes, or systems of colliding atoms as in Bose–Einstein condensate. According to Merriam-Webster and the Online Etymology Dictionary, the word "molecule" derives from the Latin "moles" or small unit of mass. Molecule – "extremely minute particle", from French molécule, from New Latin molecula, diminutive of Latin moles "mass, barrier". A vague meaning at first; the definition of the molecule has evolved. Earlier definitions were less precise, defining molecules as the smallest particles of pure chemical substances that still retain their composition and chemical properties; this definition breaks down since many substances in ordinary experience, such as rocks and metals, are composed of large crystalline networks of chemically bonded atoms or ions, but are not made of discrete molecules.
Molecules are held together by ionic bonding. Several types of non-metal elements exist only as molecules in the environment. For example, hydrogen only exists as hydrogen molecule. A molecule of a compound is made out of two or more elements. A covalent bond is a chemical bond; these electron pairs are termed shared pairs or bonding pairs, the stable balance of attractive and repulsive forces between atoms, when they share electrons, is termed covalent bonding. Ionic bonding is a type of chemical bond that involves the electrostatic attraction between oppositely charged ions, is the primary interaction occurring in ionic compounds; the ions are atoms that have lost one or more electrons and atoms that have gained one or more electrons. This transfer of electrons is termed electrovalence in contrast to covalence. In the simplest case, the cation is a metal atom and the anion is a nonmetal atom, but these ions can be of a more complicated nature, e.g. molecular ions like NH4+ or SO42−. An ionic bond is the transfer of electrons from a metal to a non-metal for both atoms to obtain a full valence shell.
Most molecules are far too small to be seen with the naked eye. DNA, a macromolecule, can reach macroscopic sizes, as can molecules of many polymers. Molecules used as building blocks for organic synthesis have a dimension of a few angstroms to several dozen Å, or around one billionth of a meter. Single molecules cannot be observed by light, but small molecules and the outlines of individual atoms may be traced in some circumstances by use of an atomic force microscope; some of the largest molecules are supermolecules. The smallest molecule is the diatomic hydrogen, with a bond length of 0.74 Å. Effective molecular radius is the size; the table of permselectivity for different substances contains examples. The chemical formula for a molecule uses one line of chemical element symbols and sometimes al
In chemistry, the valence or valency of an element is a measure of its combining power with other atoms when it forms chemical compounds or molecules. The concept of valence developed in the second half of the 19th century and helped explain the molecular structure of inorganic and organic compounds; the quest for the underlying causes of valence led to the modern theories of chemical bonding, including the cubical atom, Lewis structures, valence bond theory, molecular orbitals, valence shell electron pair repulsion theory, all of the advanced methods of quantum chemistry. The combining power, or affinity of an atom of a given element is determined by the number of hydrogen atoms that it combines with. In methane, carbon has a valence of 4. Chlorine, as it has a valence of one, can be substituted for hydrogen, so phosphorus has a valence of 5 in phosphorus pentachloride, PCl5. Valence diagrams of a compound represent the connectivity of the elements, with lines drawn between two elements, sometimes called bonds, representing a saturated valency for each element.
The two tables below show some examples of different compounds, their valence diagrams, the valences for each element of the compound. Valence only describes connectivity. A line between atoms does not represent a pair of electrons. Valence is defined by the IUPAC as: The maximum number of univalent atoms that may combine with an atom of the element under consideration, or with a fragment, or for which an atom of this element can be substituted. An alternative modern description is: The number of hydrogen atoms that can combine with an element in a binary hydride or twice the number of oxygen atoms combining with an element in its oxide or oxides; this definition differs from the IUPAC definition as an element can be said to have more than one valence. The etymology of the words valence and valency traces back to 1425, meaning "extract, preparation", from Latin valentia "strength, capacity", from the earlier valor "worth, value", the chemical meaning referring to the "combining power of an element" is recorded from 1884, from German Valenz.
In 1789, William Higgins published views on what he called combinations of "ultimate" particles, which foreshadowed the concept of valency bonds. If, for example, according to Higgins, the force between the ultimate particle of oxygen and the ultimate particle of nitrogen were 6 the strength of the force would be divided accordingly, for the other combinations of ultimate particles; the exact inception, however, of the theory of chemical valencies can be traced to an 1852 paper by Edward Frankland, in which he combined the older theories of free radicals with thoughts on chemical affinity to show that certain elements have the tendency to combine with other elements to form compounds containing 3, i.e. in the 3-atom groups or 5, i.e. in the 5-atom groups, equivalents of the attached elements. According to him, this is the manner in which their affinities are best satisfied, by following these examples and postulates, he declares how obvious it is that This “combining power” was afterwards called quantivalence or valency.
In 1857 August Kekulé proposed fixed valences for many elements, such as 4 for carbon, used them to propose structural formulas for many organic molecules, which are still accepted today. Most 19th-century chemists defined the valence of an element as the number of its bonds without distinguishing different types of valence or of bond. However, in 1893 Alfred Werner described transition metal coordination complexes such as Cl3, in which he distinguished principal and subsidiary valences, corresponding to the modern concepts of oxidation state and coordination number respectively. For main-group elements, in 1904 Richard Abegg considered positive and negative valences, proposed Abegg's rule to the effect that their difference is 8; the Rutherford model of the nuclear atom showed that the exterior of an atom is occupied by electrons, which suggests that electrons are responsible for the interaction of atoms and the formation of chemical bonds. In 1916, Gilbert N. Lewis explained valence and chemical bonding in terms of a tendency of atoms to achieve a stable octet of 8 valence-shell electrons.
According to Lewis, covalent bonding leads to octets by the sharing of electrons, ionic bonding leads to octets by the transfer of electrons from one atom to the other. The term covalence is attributed to Irving Langmuir, who stated in 1919 that "the number of pairs of electrons which any given atom shares with the adjacent atoms is called the covalence of that atom"; the prefix co - means "together". Subsequent to that, it is now more common to speak of covalent bonds rather than valence, which has fallen out of use in higher-level work from the advances in the theory of chemical bonding, but it is still used in elementary studies, where it provides a heuristic introduction to the subject. In the 1930s, Linus Pauling proposed that there are polar covalent bonds, which are intermediate between covalent and ionic, that the degree of ionic character depends on the difference of electronegativity of the two bonded atoms. Pauling also
London dispersion force
London dispersion forces are a type of force acting between atoms and molecules. They are part of the van der Waals forces; the LDF is named after the German-American physicist Fritz London. London forces are exhibited by all molecules; the electron distribution around an atom or molecule undergoes fluctuations in time. These fluctuations create instantaneous electric fields which are felt by other nearby atoms and molecules, which in turn adjust the spatial distribution of their own electrons; the net effect is that the fluctuations in electron positions in one atom induce a corresponding redistribution of electrons in other atoms, such that the electron motions become correlated. While the detailed theory requires a quantum-mechanical explanation, the effect is described as the formation of instantaneous dipoles that attract each other; the magnitude of the London dispersion force is described in terms of a single parameter called the Hamaker constant symbolized A. For atoms that are located closer together than the wavelength of light, the interaction is instantaneous and is described in terms of a "non-retarded" Hamaker constant.
For entities that are farther apart, the finite time required for the fluctuation at one atom to be felt at a second atom requires use of a "Retarded" Hamaker constant. While the London dispersion force between individual atoms and molecules is quite weak and decreases with separation like 1 R 6, in condensed matter, the effect is cumulative over the volume of materials,or within and between organic molecules, such that London dispersion forces can be quite strong in bulk solid and liquids and decays much more with distance. For example, the total force per unit area between two bulk solids decreases like 1 R 3 where R is the separation between them; the effects of London dispersion forces are most obvious in systems that are non-polar, such as hydrocarbons and symmetric molecules such as bromine, iodine. In hydrocarbons and waxes the dispersion forces are sufficient to cause condensation from the gas phase into the liquid or solid phase. Sublimation heats of e.g. hydrocarbon crystals reflect the dispersion interaction.
Liquification of oxygen and nitrogen gases into liquid phases is dominated by attractive London dispersion forces. When atoms/molecules are separated by a third medium, the situation becomes more complex. In aqueous solutions, the effects of dispersion forces between atoms/molecules are less pronounced due to competition with polarizable solvent molecules; that is, the instantaneous fluctuations in one atom or molecule are felt both by the solvent and by other molecules. London forces become stronger as the atom in question becomes larger, to a smaller degree for large molecules; this is due to the increased polarizability of molecules with larger, more dispersed electron clouds. The polarizability is a measure of ease; this trend is exemplified by the halogens. The same increase of dispersive attraction occurs within and between organic molecules in the order RF<RCL<RBr<RI, or with other more polarizable heteroatoms. Fluorine and chlorine are gases at room temperature, bromine is a liquid, iodine is a solid.
The London forces are thought to be arise from the motion of electrons. Thus the forces of attraction b/w; the first explanation of the attraction between noble gas atoms was given by Fritz London in 1930. He used a quantum-mechanical theory based on second-order perturbation theory; the perturbation is because of the Coulomb interaction between the electrons and nuclei of the two moieties. The second-order perturbation expression of the interaction energy contains a sum over states; the states appearing in this sum are simple products of the stimulated electronic states of the monomers. Thus, no intermolecular antisymmetrization of the electronic states is included and the Pauli exclusion principle is only satisfied. London wrote Taylor series expansion of the perturbation in 1 R, where R is the distance between the nuclear centers of mass of the moieties; this expansion is known as the multipole expansion because the terms in this series can be regarded as energies of two interacting multipoles, one on each monomer.
Substitution of the multipole-expanded form of V into the second-order energy yields an expression that resembles somewhat an expression describing the interaction between instantaneous multipoles. Additionally, an approximation, named after Albrecht Unsöld, must be introduced in order to obtain a description of London dispersion in terms of dipole polarizabilities and ionization potentials. In this manner, the following approximation is obtained for the dispersion interaction E A B d i s p between two atoms A and
Vapor pressure or equilibrium vapor pressure is defined as the pressure exerted by a vapor in thermodynamic equilibrium with its condensed phases at a given temperature in a closed system. The equilibrium vapor pressure is an indication of a liquid's evaporation rate, it relates to the tendency of particles to escape from the liquid. A substance with a high vapor pressure at normal temperatures is referred to as volatile; the pressure exhibited by vapor present above a liquid surface is known as vapor pressure. As the temperature of a liquid increases, the kinetic energy of its molecules increases; as the kinetic energy of the molecules increases, the number of molecules transitioning into a vapor increases, thereby increasing the vapor pressure. The vapor pressure of any substance increases non-linearly with temperature according to the Clausius–Clapeyron relation; the atmospheric pressure boiling point of a liquid is the temperature at which the vapor pressure equals the ambient atmospheric pressure.
With any incremental increase in that temperature, the vapor pressure becomes sufficient to overcome atmospheric pressure and lift the liquid to form vapor bubbles inside the bulk of the substance. Bubble formation deeper in the liquid requires a higher temperature due to the higher fluid pressure, because fluid pressure increases above the atmospheric pressure as the depth increases. More important at shallow depths is the higher temperature required to start bubble formation; the surface tension of the bubble wall leads to an overpressure in the small, initial bubbles. Thus, thermometer calibration should not rely on the temperature in boiling water; the vapor pressure that a single component in a mixture contributes to the total pressure in the system is called partial pressure. For example, air at sea level, saturated with water vapor at 20 °C, has partial pressures of about 2.3 kPa of water, 78 kPa of nitrogen, 21 kPa of oxygen and 0.9 kPa of argon, totaling 102.2 kPa, making the basis for standard atmospheric pressure.
Vapor pressure is measured in the standard units of pressure. The International System of Units recognizes pressure as a derived unit with the dimension of force per area and designates the pascal as its standard unit. One pascal is one newton per square meter. Experimental measurement of vapor pressure is a simple procedure for common pressures between 1 and 200 kPa. Most accurate results are obtained near the boiling point of substances and large errors result for measurements smaller than 1kPa. Procedures consist of purifying the test substance, isolating it in a container, evacuating any foreign gas measuring the equilibrium pressure of the gaseous phase of the substance in the container at different temperatures. Better accuracy is achieved when care is taken to ensure that the entire substance and its vapor are at the prescribed temperature; this is done, as with the use of an isoteniscope, by submerging the containment area in a liquid bath. Low vapor pressures of solids can be measured using the Knudsen effusion cell method.
In a medical context, vapor pressure is sometimes expressed in other units millimeters of mercury. This is important for volatile anesthetics, most of which are liquids at body temperature, but with a high vapor pressure. Anesthetics with a higher vapor pressure at body temperature will be excreted more as they are exhaled from the lungs; the Antoine equation is a mathematical expression of the relation between the vapor pressure and the temperature of pure liquid or solid substances. The basic form of the equation is: log P = A − B C + T and it can be transformed into this temperature-explicit form: T = B A − log P − C where: P is the absolute vapor pressure of a substance T is the temperature of the substance A, B and C are substance-specific coefficients log is either log 10 or log e A simpler form of the equation with only two coefficients is sometimes used: log P = A − B T which can be transformed to: T = B A − log P Sublimations and vaporizations of the same substance have separate sets of Antoine coefficients, as do components in mixtures.
Each parameter set for a specific compound is only applicable over a specified temperature range. Temperature ranges are chosen to maintain the equation's accuracy of a few up to 8–10 percent. For many volatile substances, several different sets of parameters are available and used for different temperature ranges; the Antoine equation has poor accuracy with any single parameter set when used from a compound's melting point to its critical temperature. Accuracy is usually poor when vapor pressure is under 10 Torr because of the limitations of the apparatus used to establish the Antoine parameter values; the Wagner equation gives "o
Hydrogen is a chemical element with symbol H and atomic number 1. With a standard atomic weight of 1.008, hydrogen is the lightest element in the periodic table. Hydrogen is the most abundant chemical substance in the Universe, constituting 75% of all baryonic mass. Non-remnant stars are composed of hydrogen in the plasma state; the most common isotope of hydrogen, termed protium, has no neutrons. The universal emergence of atomic hydrogen first occurred during the recombination epoch. At standard temperature and pressure, hydrogen is a colorless, tasteless, non-toxic, nonmetallic combustible diatomic gas with the molecular formula H2. Since hydrogen forms covalent compounds with most nonmetallic elements, most of the hydrogen on Earth exists in molecular forms such as water or organic compounds. Hydrogen plays a important role in acid–base reactions because most acid-base reactions involve the exchange of protons between soluble molecules. In ionic compounds, hydrogen can take the form of a negative charge when it is known as a hydride, or as a positively charged species denoted by the symbol H+.
The hydrogen cation is written as though composed of a bare proton, but in reality, hydrogen cations in ionic compounds are always more complex. As the only neutral atom for which the Schrödinger equation can be solved analytically, study of the energetics and bonding of the hydrogen atom has played a key role in the development of quantum mechanics. Hydrogen gas was first artificially produced in the early 16th century by the reaction of acids on metals. In 1766–81, Henry Cavendish was the first to recognize that hydrogen gas was a discrete substance, that it produces water when burned, the property for which it was named: in Greek, hydrogen means "water-former". Industrial production is from steam reforming natural gas, less from more energy-intensive methods such as the electrolysis of water. Most hydrogen is used near the site of its production, the two largest uses being fossil fuel processing and ammonia production for the fertilizer market. Hydrogen is a concern in metallurgy as it can embrittle many metals, complicating the design of pipelines and storage tanks.
Hydrogen gas is flammable and will burn in air at a wide range of concentrations between 4% and 75% by volume. The enthalpy of combustion is −286 kJ/mol: 2 H2 + O2 → 2 H2O + 572 kJ Hydrogen gas forms explosive mixtures with air in concentrations from 4–74% and with chlorine at 5–95%; the explosive reactions may be triggered by heat, or sunlight. The hydrogen autoignition temperature, the temperature of spontaneous ignition in air, is 500 °C. Pure hydrogen-oxygen flames emit ultraviolet light and with high oxygen mix are nearly invisible to the naked eye, as illustrated by the faint plume of the Space Shuttle Main Engine, compared to the visible plume of a Space Shuttle Solid Rocket Booster, which uses an ammonium perchlorate composite; the detection of a burning hydrogen leak may require a flame detector. Hydrogen flames in other conditions are blue; the destruction of the Hindenburg airship was a notorious example of hydrogen combustion and the cause is still debated. The visible orange flames in that incident were the result of a rich mixture of hydrogen to oxygen combined with carbon compounds from the airship skin.
H2 reacts with every oxidizing element. Hydrogen can react spontaneously and violently at room temperature with chlorine and fluorine to form the corresponding hydrogen halides, hydrogen chloride and hydrogen fluoride, which are potentially dangerous acids; the ground state energy level of the electron in a hydrogen atom is −13.6 eV, equivalent to an ultraviolet photon of 91 nm wavelength. The energy levels of hydrogen can be calculated accurately using the Bohr model of the atom, which conceptualizes the electron as "orbiting" the proton in analogy to the Earth's orbit of the Sun. However, the atomic electron and proton are held together by electromagnetic force, while planets and celestial objects are held by gravity; because of the discretization of angular momentum postulated in early quantum mechanics by Bohr, the electron in the Bohr model can only occupy certain allowed distances from the proton, therefore only certain allowed energies. A more accurate description of the hydrogen atom comes from a purely quantum mechanical treatment that uses the Schrödinger equation, Dirac equation or the Feynman path integral formulation to calculate the probability density of the electron around the proton.
The most complicated treatments allow for the small effects of special relativity and vacuum polarization. In the quantum mechanical treatment, the electron in a ground state hydrogen atom has no angular momentum at all—illustrating how the "planetary orbit" differs from electron motion. There exist two different spin isomers of hydrogen diatomic molecules that differ by the relative spin of their nuclei. In the orthohydrogen form, the spins of the two protons are parallel and form a triplet state with a molecular spin quantum number of 1. At standard temperature and pressure, hydrogen gas contains about 25% of the para form and 75% of the ortho form known as the "normal form"; the equilibrium ratio of orthohydrogen to parahydrogen depends on temperature, but because the ortho form is an excited state and has a higher energy
Oxygen is the chemical element with the symbol O and atomic number 8. It is a member of the chalcogen group on the periodic table, a reactive nonmetal, an oxidizing agent that forms oxides with most elements as well as with other compounds. By mass, oxygen is the third-most abundant element in the universe, after helium. At standard temperature and pressure, two atoms of the element bind to form dioxygen, a colorless and odorless diatomic gas with the formula O2. Diatomic oxygen gas constitutes 20.8% of the Earth's atmosphere. As compounds including oxides, the element makes up half of the Earth's crust. Dioxygen is used in cellular respiration and many major classes of organic molecules in living organisms contain oxygen, such as proteins, nucleic acids and fats, as do the major constituent inorganic compounds of animal shells and bone. Most of the mass of living organisms is oxygen as a component of water, the major constituent of lifeforms. Oxygen is continuously replenished in Earth's atmosphere by photosynthesis, which uses the energy of sunlight to produce oxygen from water and carbon dioxide.
Oxygen is too chemically reactive to remain a free element in air without being continuously replenished by the photosynthetic action of living organisms. Another form of oxygen, ozone absorbs ultraviolet UVB radiation and the high-altitude ozone layer helps protect the biosphere from ultraviolet radiation. However, ozone present at the surface is a byproduct of thus a pollutant. Oxygen was isolated by Michael Sendivogius before 1604, but it is believed that the element was discovered independently by Carl Wilhelm Scheele, in Uppsala, in 1773 or earlier, Joseph Priestley in Wiltshire, in 1774. Priority is given for Priestley because his work was published first. Priestley, called oxygen "dephlogisticated air", did not recognize it as a chemical element; the name oxygen was coined in 1777 by Antoine Lavoisier, who first recognized oxygen as a chemical element and characterized the role it plays in combustion. Common uses of oxygen include production of steel and textiles, brazing and cutting of steels and other metals, rocket propellant, oxygen therapy, life support systems in aircraft, submarines and diving.
One of the first known experiments on the relationship between combustion and air was conducted by the 2nd century BCE Greek writer on mechanics, Philo of Byzantium. In his work Pneumatica, Philo observed that inverting a vessel over a burning candle and surrounding the vessel's neck with water resulted in some water rising into the neck. Philo incorrectly surmised that parts of the air in the vessel were converted into the classical element fire and thus were able to escape through pores in the glass. Many centuries Leonardo da Vinci built on Philo's work by observing that a portion of air is consumed during combustion and respiration. In the late 17th century, Robert Boyle proved. English chemist John Mayow refined this work by showing that fire requires only a part of air that he called spiritus nitroaereus. In one experiment, he found that placing either a mouse or a lit candle in a closed container over water caused the water to rise and replace one-fourteenth of the air's volume before extinguishing the subjects.
From this he surmised that nitroaereus is consumed in both combustion. Mayow observed that antimony increased in weight when heated, inferred that the nitroaereus must have combined with it, he thought that the lungs separate nitroaereus from air and pass it into the blood and that animal heat and muscle movement result from the reaction of nitroaereus with certain substances in the body. Accounts of these and other experiments and ideas were published in 1668 in his work Tractatus duo in the tract "De respiratione". Robert Hooke, Ole Borch, Mikhail Lomonosov, Pierre Bayen all produced oxygen in experiments in the 17th and the 18th century but none of them recognized it as a chemical element; this may have been in part due to the prevalence of the philosophy of combustion and corrosion called the phlogiston theory, the favored explanation of those processes. Established in 1667 by the German alchemist J. J. Becher, modified by the chemist Georg Ernst Stahl by 1731, phlogiston theory stated that all combustible materials were made of two parts.
One part, called phlogiston, was given off when the substance containing it was burned, while the dephlogisticated part was thought to be its true form, or calx. Combustible materials that leave little residue, such as wood or coal, were thought to be made of phlogiston. Air did not play a role in phlogiston theory, nor were any initial quantitative experiments conducted to test the idea. Polish alchemist and physician Michael Sendivogius in his work De Lapide Philosophorum Tractatus duodecim e naturae fonte et manuali experientia depromti described a substance contained in air, referring to it as'cibus vitae', this substance is identical with oxygen. Sendivogius, during his experiments performed between 1598 and 1604, properly recognized that the substance is equivalent to the gaseous byproduct released by the thermal decomposition of potassium nitrate. In Bugaj’s view, the isolation of oxygen and the proper association of the substance to that part of air, required for life, lends sufficient weight to the discovery of oxygen by Sendivogius.